1. Chapter 3
1.1. Light
1.1.1. form of electromagnetic radiation
1.1.2. Travel as wave
1.1.3. Speed of Light- 3.00 x 10^8 m/s
1.2. Wave
1.2.1. periodic oscillation that transmits energy through space
1.2.2. Wavelength
1.2.2.1. distance from one peak to the next
1.2.3. Amplitude
1.2.3.1. height from the centerline of wave peak to the next
1.2.4. Frequency
1.2.4.1. How fast wave peaks are going
1.2.5. EM Waves
1.2.5.1. Interaction between waves is interface
1.3. Photons
1.3.1. Electromagnetic radiation is a stream of “particles”
1.3.2. h - Planck's constant; v- radiation frequency; upside down Y is radiation wavelength
1.3.3. Ephoton= hv=hc/ (upsidedown) Y
1.4. Photoelectric Effect
1.4.1. minimum energy required to remove electron Eo=hvo
1.4.2. Low frequency= no electrons emitting
1.4.3. Einstein's Idea: Light is emitted!!
1.4.3.1. One photon→ one electron emitted
1.5. Emission
1.5.1. light emitted from excited atom
1.5.2. Occurs when electrons falls to lower level
1.6. Absorption
1.6.1. missing wavelength when white light passes through sample
1.6.2. Occurs when light of certain wavelength causes electrons to jump to higher level
2. Chapter 3 cont.
2.1. Electron Affiities
2.1.1. Energy released when a neutral atom in the gas phase gains an electron
2.1.1.1. The more negative, the greater tendency to accept an electron/ electron affinity
2.1.1.1.1. Decreases (less negative) going down a group
2.1.1.1.2. Increases (more negative) going across period
2.1.1.1.3. Exceptions: F & O out of line
2.2. Trends
2.2.1. Larger Zeff, more energy required to remove an electron
2.2.2. Larger n, less energy required
2.2.3. Closer to NG, harder to remove
2.2.4. 1st IE- increases across period
2.3. Ionization
2.3.1. Ionization energy
2.3.1.1. energy required to move electrons from the ground state of neutral atom in gas phase
2.4. Electron Configuration
2.4.1. Valence Electrons- Electrons in all sublevels with highest energy shell
2.4.2. Core electrons- Electrons in lower energy shell
2.4.3. With Ions
2.4.3.1. Results from gain/loss of electrons
2.4.3.2. Cations form with Main Group Metals
2.4.3.3. Anions form with Non-Metals
2.4.3.3.1. Fill Valence orbitals
2.4.3.3.2. Group # - 8 give charge
2.5. Quantum Numbers
2.5.1. Explain importance of concepts of energy levels, orbitals, orbital size, orbital orientation, and electric spin
2.5.2. Identify s,p,d orbitals
2.5.3. Principal quantum # (n): related to the size & energy of the orbital
2.5.4. Angular momentum (l): related to shape of atomic orbitals (subshell)
2.5.4.1. S orbitals; l=0 One orbital l=(n-1) P orbitals; l=1 Three orbitals D orbitals; l=2 Five orbitals F orbitals; l=3 Seven orbitals
3. Chapter 5
3.1. Molecular Shape
3.1.1. Bond Angle
3.1.1.1. Angle (in degrees) defined by lines joining the centers of two atoms to the center of a third atom to which they are covalently bonded
3.1.2. VSEPR Model
3.1.2.1. Valence Shell Electron Pair Repulsion Model
3.1.3. Based on regions of high electron density around the central atom
3.1.4. Refers to the arrangement of the different atoms around the central atom in a molecule or ion
3.2. Electron Group Geometry
3.2.1. Linear
3.2.1.1. Two electron groups
3.2.1.2. Two bonds and 0 lone pairs
3.2.1.3. Bond angle is 180 degrees
3.2.2. Trigonal Planar
3.2.2.1. Three electron groups
3.2.2.2. 3 bonds and 0 lone pairs
3.2.2.3. 120 degrees
3.2.2.4. Bent Molecular Shape
3.2.2.4.1. Bond angle < 120 degrees
3.2.3. Tetrahedral
3.2.3.1. Four charge regions
3.2.3.2. Bond angle is 109.5 degrees
3.2.3.3. 4 bonds and 0 lone pairs
3.2.3.4. Tetrahedral Molecular Shape
3.2.3.4.1. Bond angle= 109.5 degrees
3.2.3.5. Trigonal Pyramidal Molecular Shape
3.2.3.5.1. Bond angle is < 109.5 degrees
3.2.3.6. Bent Molecular Shape
3.2.3.6.1. Bond angle < 109.5 degrees
3.2.4. Bipyramidal
3.2.4.1. 5 Charge Regions
3.2.4.2. TWO LONE PAIR: T_SHAPED MOLECULAR SHAPE
3.2.4.3. ONE LONE PAIR: SEE SAW MOLECULAR SHAPE
3.2.4.4. No lone pair; Same as trigonal bipyramidal molecular shape
3.2.5. Octahedral
3.2.5.1. 6 charge regions
3.2.5.2. Bond angles = 90 degrees
3.2.6. Refers to the arrangement of electron pairs
3.3. Atomic Orbital
3.3.1. Wave function whose square gives the probability of finding an electron within a given region of space in an atom
3.4. Molecular Orbital
3.4.1. Wave function whose square gives the probability of finding an electron within a given region of space in an atom
3.4.1.1. Specific energy & shape, contain a maximum of two electrons, described by wave functions, electron density distribution
3.4.2. Bonding Orbitals
3.4.2.1. Region of increased electron density between nuclear centers that hold atoms together
3.4.3. Anti bonding Orbitals
3.4.3.1. Regions of electron density that destabilize the molecule because they do not increase electron density between nuclear center
3.4.4. MO Diagram
3.4.4.1. Energy level diagram for showing the relative energy goes and occupancy of the MOs for a molecule
3.4.5. Sigma Bond
3.4.5.1. Covalent bond with the highest electron density along the bond axis
3.4.5.1.1. Head to head overlap of orbitals
3.4.5.1.2. Strongest type of covalent
3.4.6. Pi Bond
3.4.6.1. Formed by mixing of atomic orbitals not oriented along the bonding axis in a molecule
3.4.6.1.1. Sideways overlap of orbitals
3.4.6.1.2. Weaker than sigma bonds
3.5. Valence Bond Theory
3.5.1. Valence electrons of the atoms in a molecule reside in quantum-mechanical atomic orbitals
3.5.2. Bonding takes place between atoms
3.6. Hybridizing
3.6.1. Mixing atomic orbitals into new hybrid orbitals
3.7. Bond Order
3.7.1. BO = ½ (# bonding electrons - # anti bonding electrons)
4. Chapter 4
4.1. Elements
4.1.1. Some exist as molecules
4.1.2. Molecular Elements
4.1.2.1. Smallest particle involves two or more atoms bonded
4.2. Molecular Compounds
4.2.1. Covalently bonded elements
4.2.2. Non-metal & nonmetal/metalloid
4.2.3. NO METALS
4.2.4. Electrons shared equally
4.3. Types of Bonds
4.3.1. Ionic
4.3.1.1. Between metal & non-metal
4.3.1.2. Attraction of opposite charges
4.3.1.3. One atom loses valence electrons & other gained electrons
4.3.1.4. Cation electron deficient; anion excess electron density
4.3.1.5. High melting point in solids
4.3.1.6. Non-directional bond
4.3.2. Polar Covalent
4.3.2.1. Between non-metal & nonmetal/metalloid
4.3.2.2. At low, there's a balance
4.3.2.3. Unequal sharing of electrons
4.3.2.4. Non-uniform charge distribution
4.3.2.5. Involves atoms of different elements
4.3.2.6. Low melting point
4.3.2.7. Directional bonds
4.3.2.8. Molecule*- a collection of atoms that are strongly attracted to one another & move & act together as a single particle
4.3.3. Pure (non-polar) Covalent
4.3.3.1. Equal sharing
4.3.3.2. Uniform distribution
4.3.3.3. Involves atoms of same elements
4.3.4. Metallic Bonds
4.3.4.1. Involves metal or metals
4.3.4.2. Good conductors
4.3.4.3. Overlap of metal atom valence orbitals
4.3.4.3.1. Valence electrons are not localized to one atom but are delocalized over many atoms
4.4. Electronegativity
4.4.1. measure of the ability of an atom to attract shared electrons in a chemical bond
4.4.2. The smaller the atom the greater its electronegativity.
4.4.3. The larger the difference in electronegativity between two atoms, the more polar the bond between the two atoms—the more unequal the electron distribution
4.4.4. Trends
4.4.4.1. Increases across a period
4.4.4.2. Decreases down a family/group
4.5. Dipole Movement
4.5.1. A measure of bond polarity
4.5.1.1. The larger the dipole moment the more polar a covalent bond
4.6. Bond Type
4.6.1. If difference is 0.0 to 0.4, the bond is non polar covalent.
4.6.2. If difference is 0.5 to 1.9, the bond is polar covalent.
4.6.3. If difference is ≥2.0, the bond is ionic
4.7. Lattice Energy
4.7.1. A measure of the electrostatic interaction energies between ions in a solid
4.7.1.1. Ionic Bonding- Ions held together by electrostatic attraction of opposite charges
4.7.1.2. Directly related to charge; indirectly related to the distance between ions
4.7.1.3. The larger q1/q2, the larger the lattice energy
4.7.1.4. Lattice energy inc→ melting point inc
4.7.1.5. Smaller the ions, the larger the lattice energy
5. Chapter 4 cont.
5.1. Polyatomic Ions
5.1.1. A group of atoms that are covalently bonded together but entire group carries a charge
5.2. Writing formulas
5.2.1. Identify cation and anion charges
5.2.1.1. Metal forms cation
5.2.1.2. Non metal forms anion
5.2.2. Cation written first
5.2.3. Adjust subscripts so total positive charge = total negative charge so neutral
5.2.4. Check for smallest whole number ratio
5.3. Naming compounds
5.3.1. First word is name of cation/ parent element
5.3.2. Second word is name of anion/ parent element with -ide
5.3.3. CHARGE MUST BE BALANCE
5.4. Lewis Dot Structures
5.4.1. Dots represent valence electrons
5.4.1.1. Bonding Pair- Two electrons are shared by atoms
5.4.1.2. Lone pair- Two electrons are not shared by atoms but belong to particular atom
5.4.1.3. Each bond shares 2 electrons
5.4.2. Shows placement of bonding & non bonding electron pairs in the molecule
5.4.3. Predicts shape
5.4.4. Chemical symbol represents nucleus and core electrons
5.4.5. How to Draw
5.4.5.1. Determine number of valence electrons for molecule
5.4.5.2. Write skeletal structure
5.4.5.3. Distribute electrons
5.4.5.4. Complete octets by making double and triple bonds
5.4.5.5. Check
5.4.6. Octet has 8 electrons
5.4.7. B,C,N,O,F never expand octet
5.5. Resonance
5.5.1. More than one valid ELectron Dot Structure can be written for a molecule
5.5.2. Total Number of bonds/number of domains
5.5.2.1. Fractions have resonance
5.5.3. Only electrons move, atoms never move
5.5.4. Lone pairs and multiple bonds move; single bonds never move
5.5.5. Number of electrons never changes; all structures have SAME net charge
5.6. Formal Charge
5.6.1. Charge an atom in a molecule or ion would have if all bonding electrons were shared equally between the bonded atoms
5.6.2. Formal Charge= # valence electrons - (free electrons+bonds)
5.6.3. Sum of all formal charges in molecule must be zero